Atoms 101: A Beginner's Guide

The ancient Greeks were not nearly as stupid as they looked, despite draping themselves in bed sheets for every occasion. It was they who proposed the idea of the atom, or atomos meaning 'indivisible'.

Back in those days, philosophers were the early scientists, and it was Leucippus who said that everything is made of either tiny atoms or voids in between them.

While this was a great starter hypothesis, today we know the atom itself is largely comprised of empty space. It's also made of even tinier quantum particles, the existence of which can only be inferred indirectly due to the insanity of the quantum world.

So the Greek name is a misnomer: atoms are indeed divisible, and what quantum quirks lurk within are profoundly unsettling.

What Do Atoms Look Like?

Tricky one, this. Neils Bohr devised the conceptual model you learned in school, which we'll look at here in the form of the element helium.

Bear in mind, this is merely a representative model, because actually trying to draw an atom is like riding a bicycle with the handlebars on backwards while counting in thirty-sevens. It's difficult.

At the centre of the atom is the nucleus, home to two different particles:

Protons are positively charged particles

Neutrons are neutral particles with no electrical charge

Orbiting the nucleus in various energy shells are another, much smaller type of particle:

Electrons are negatively charged particles

These features alone tell us a lot about the species of atom (carbon, oxygen, hydrogen, and so on) and its instinctive behaviours (the types of bonds it forms with other atoms).

The Atomic Staircase

Now look at this table.

We distinguish elements from one another by their number of protons. This, combined with their number of neutrons, have implications for their atomic mass. Important figures, these.

Imagine a staircase, where each step represents the increasing number of protons in the various elements. Let's climb.

On the first step, the lightest atom is hydrogen with one proton. Hydrogen is unique in that it has zero neutrons, so it has an atomic mass of just one. It has one electron too, but let's not worry about electrons at this point.

On the second step, we greet helium. It has two protons and two neutrons, which add up to an atomic mass of four.

The third step features lithium, with three protons and four neutrons, giving it an atomic mass of seven. You get the idea.

In nature, the staircase has 92 steps, finishing around uranium. However, physicists have figured out how to make new elements, thereby adding an extra 26 steps to the staircase, for a total of 118 elements.

All man-made elements are radioactive and many are named for historical scientists, such as the awkward-sounding rutherfordium and einsteinium.

The Electron Orbits

Now consider the tiny negative electrons which sit in the orbits of these atoms. These play a crucial role in determining how atoms react to each another, because it's the orbits that interact physically.

While the protons in the nucleus define the positive electrical charge, the electrons in the orbits balance them out with negative charge.

Although electrons are relatively minuscule, each one still holds an equal and opposite charge to one proton. This makes the entire atom electrically neutral in its default state.

Atomic Bonds

Like human beings, atoms tend to be very social creatures, and like to bond with one another when the opportunity arises. This is how we get the molecules and compounds necessary to all life on Earth.

Now we need to add some seemingly unintuitive terminology to your chemistry vocab. But don't worry, there's only three of them, which you'll agree is 33% more manageable than four*.

*Feel free to disagree.

Bond #1. Covalent bonds involve sharing electrons

This is the most common way to bond. Two atoms team-up and share pairs of electrons in their outer orbits.

There are basic mathematical rules to electron orbits. The first, innermost orbit is deemed full when it contains two electrons. Many orbits thereafter are full once they contain eight electrons.

This apparent simplicity is what drives the behaviour of many atoms to bond with others.

Take oxygen for example.

Oxygen - with a total of eight electrons - has a problem. Two of its electrons are happily contained in its inner orbit, but its outer orbit hosts only six of eight potential seats at the dinner table.

This sucks for oxygen, who dreams of having two more guests to dinner.

The easiest way do this is to invite two electrons from the outer orbits of other atoms and share them with their host atoms.

This is how we get a water molecule. An oxygen atom forms two covalent bonds with two hydrogen atoms to form H2O.

Note that hydrogen has the same dilemma, seeking one more electron to fulfil its one and only orbit. Hydrogen and oxygen are quite naturally drawn to each other for this reason.

Water is one of the universe's best ideas so far. It makes up 90% of your body and is the medium for the chemical reactions in your cells.

It also makes up plant sap, which is a huge deal because plants are the original source of food for all animals on the planet.

Bond #2. Ionic bonds involve stealing electrons

Another, less common, form of atom joinery is called ionic bonding.

Instead of politely sharing electrons, some atoms outright steal an electron from another atom altogether.

This is fine, though, because their victims are actually willing participants. There's no foul play in physics.

To make table salt, for instance, a sodium atom gives up the one and only electron in its outer orbit to an atom of chlorine.

Sodium is perfectly happy with this arrangement because donating an electron enables it to drop the frivolous orbit entirely. Its next innermost orbit is full, and so sodium can relax for a bit.

Chlorine is in the opposite position because it needs to gain just one electron to complete its outer shell.

It could form a covalent bond in this situation, but if sodium's lurking around, it takes advantage by forming this ionic bond.

There's a key difference between ionic and covalent bonds. Note how sodium is left positively charged by losing an electron, and chlorine is left negatively charged by gaining one.

When an atom gains or loses an electron like this, it becomes an ion.

And it's the opposite electrical charges that now attract the ions to each other, as opposed to the continued requirement to share an electron pair.

What Are Isotopes?

We've looked at what happens when you gain or lose protons and electrons. So what happens if you gain or lose a neutron?

This is where your atom becomes an isotope. The atomic mass has changed and this is reflected in its new name.

Look at carbon and its isotopes.

In the atmosphere, cosmic rays smack into carbon atoms and provide them with extra neutrons. Carbon can go from having six neutrons to seven or even eight neutrons. It's now a carbon isotope.

The standard carbon atom becomes a carbon-13 or carbon-14 isotope, respectively, where the number reflects the abnormal atomic weight.

The Quantum World

The Bohr model of atoms suggests electrons orbit the nucleus of an atom in the same way the planets orbit the sun. Not only is this analogy over-simplistic, it's also completely wrong. Sorry about that.

Electrons more accurately pop in and out of existence somewhere in their designated energy shells.

This behaviour can be predicted, in theory. Quantum physics tells us that electron orbits (better described as three-dimensional energy shells) are a mathematical function. They describe the probability of electrons being present at particular locations at any particular moment.

To visualise this effect, the Deviant Artist DarkSilverflame created a program to plot the probability of where electrons can be found in hydrogen atoms.